Question

The activation energy of a reaction is zero. Its rate constant at $280 \,K$ is $1.6 \times 10^{-6} s^{-1}$ the rate constant at $300\, K$ is

• A. $3.2\times 10^{-6} \,s^{-1}$
• B. zero
• C. $1.6\times 10^{-6} \,s^{-1}$
• D. $1.6\times 10^{-5} \,s^{-1}$

Answer 1

Answer: (C)

Key Idea Arrhenius equation is given as:

$\log \frac{k_{2}}{k_{1}}=\frac{E_{a}}{2.303 R}\left[\frac{1}{T_{1}}-\frac{1}{T_{2}}\right]$

Given,

Activation energy of a reaction, $E_{A}=0$

Rate constant, $k_{1}=1.6 \times 10^{-6} s ^{-1}$

Temperature, $T_{1}=280 K , T_{2}=300 \,K$

According to Arrhenius equation

$log \frac{k_{2}}{1.6 \times 10^{-6}}=\frac{0}{2.303 R}\left[\frac{1}{280}-\frac{1}{300}\right]$

$log \frac{k_{2}}{1.6 \times 10^{-6}}=0$

$\frac{k_{2}}{1.6 \times 10^{-6}}=$ antilog 0

$\frac{k_{2}}{1.6 \times 10^{-6}}=1$

$\therefore k_{2}=1.6 \times 10^{-6} \,s ^{-1}$